Week 2 F11 M - Atomic
size
 • 

Info iconThis preview shows page 1. Sign up to view the full content.

View Full Document Right Arrow Icon
This is the end of the preview. Sign up to access the rest of the document.

Unformatted text preview: Atomic
size
 •  As
the
nuclear
charge
increases,
the
electrons
are
 pulled
in
toward
the
center
of
the
atom,
and
the
 size
of
any
par8cular
orbital
decreases.
 •  On
the
other
hand,
as
the
nuclear
charge
 increases,
more
electrons
are
added
to
the
atom,
 and
their
mutual
repulsion
keeps
the
outer
 orbitals
large.
 •  The
interac8on
of
the
two
effects
results
in
 gradual
decrease
in
atomic
size
across
a
period.
 Shielding
effect.
Slater
Rule.
 Varia8on
in
Size
of
Atoms
 Rela8ve
Size
of
Select
Ions
and
Their
 Parent
Atoms
 Ioniza8on
Energy
 •  First
Ioniza8on
energy
(IE1)
=
Energy
change
at
 0K
associated
with
the
removal
of
the
first
 valence
electron.
Units
(kJmol‐1
or
eV).
 •  1eV
≈
96.5
kJ/mol.
 Ioniza8on
Energy
 • 
High
values
of
IE1
associated
with
noble
gas.
 • 
Low
values
of
IE1
associated
with
Group
1.
 • 
General
increase
in
values
of
IE1
as
a
given
period
is
crossed.
(Zeff
 increases)
 • 
Discon8nuity
in
IE1
on
going
from
Element
in
group
15
to
16.
 • 
Decrease
in
values
of
IE1
on
going
from
element
in
group
2
or
12
to
13.
 Pauling’s
Electronega8vity
Values
 Minus
the
Energy
change
for
the
gain
of
an
electron
by
a
gaseous
atom.
 Homework
assignment
Chapter
1
 



1.1;
1.6;
1.7;
1.8;
1.11;
1.12;
1.13;
1.16;
1.17;
1.18;
 1.19;
1.21;
1.22;
1.23;
1.24;
1.26;
1.28;
1.30;
1.33.

 Chapter
2
 Basic
Concepts
 Molecules
 Lewis
structures
 •  The
dot
approach
to
represent
the
number
of
valence
 electrons.
 •  What
the
structures
of
the
following
molecules???
 What
is
the
atom
connec8vity???

 –  H2O,
O2,
XeF4,
PCl5
and
SF6.
 How
to
determine
Lewis
structures??
 •  Recommended
procedure:
 –  Determine
which
atom
in
the
central
atom
(single
atom,
least
 electronega8f…)
 –  Count
the
total
number
of
valence
electrons
(N) needed to account for the atoms (based on the column of the atom in the periodic table) and charge (add one electrons for each nega7ve charge; subtract one electron for each posi7ve charge). –  Draw the framework with single bonds. Some knowledge of the way the atoms are connected may be required. –  Using lone pairs, complete octets around the noncentral atoms. –  Count the number of electrons depicted (two for each bond and two for each lone pair). –  If this number is less than N, then add electrons to the central atom un7l the total number of electrons depicted is N. –  If the octet rule is not sa7sfied for the central atom and lone‐pair electrons are nearby, use those electrons to make double or triple bonds to the central atom. –  Check each atom to see if it has a formal charge. (Singly bonded oxygen will require a nega<ve charge, for example.)
 •  Resonance:
 –  there
is
more
than
one
possible
way
in
which
the
 valence
electrons
can
be
placed
in
a
Lewis
 structure.
 •  When
a
molecule
has
several
resonance
 structures,
its
overall
electronic
energy
is
 lowered
making
it
more
stable.
 Expanded
Shells
 •  When
it
is
impossible
to
draw
a
structure
consistent
 with
the
octet
rule,
it
is
necessary
to
increase
the
 number
of
electrons
around
the
central
atom.
 •  The
increased
number
of
electrons
is
described
as
an
 expanded
shell
or
an
expended
electron
count.

 •  Term
Hypervalent
is
some8mes
used
to
describe
 central
atoms
that
have
electron
counts
greater
than
 the
atom’s
usual
requirement.

 Formal
Charges
 •  Formal
charges:
 –  apparent
electronic
charge
of
each
atom
in
a
 molecule,
based
on
the
electron‐dot
structure.
 –  It
is
a
tool
for
assessing
Lewis
structure
and
NOT
a
 measure
of
any
actual
charge
on
the
atoms.

 Formal
Charge
 Charge
on
molecule
or
ion
=
sum
of
all
the
formal
charges
 Place
nega8ve
formal
charges
on
more
electronega8ve
elements
 Example
 •  •  •  •  •  •  NO2- Step one: –  Nitrogen is the least electronegative atom, so it is the central atom by multiple criteria. Step two: –  Count valence electrons. Nitrogen has 5 valence electrons; each oxygen has 6, for a total of (6 × 2) + 5 = 17. The ion has a charge of −1, which indicates an extra electron, so the total number of electrons is 18. Step three: –  Place ion pairs. Each oxygen must be bonded to the nitrogen, which uses four electrons — two in each bond. The 14 remaining electrons should initially be placed as 7 lone pairs. Each oxygen may take a maximum of 3 lone pairs, giving each oxygen 8 electrons including the bonding pair. The seventh lone pair must be placed on the nitrogen atom. Step four: Satisfy the octet rule. Both oxygen atoms currently have 8 electrons assigned to them. The nitrogen atom has only 6 electrons assigned to it. One of the lone pairs on an oxygen atom must form a double bond, but either atom will work equally well. We therefore must have a resonance structure. Step five: Tie up loose ends. Two Lewis structures must be drawn: one with each oxygen atom double-bonded to the nitrogen atom. The second oxygen atom in each structure will be singlebonded to the nitrogen atom. Place brackets around each structure, and add the charge (−) to the upper right outside the brackets. Draw a double-headed arrow between the two resonance form
 Examples
 •  Determine
the
Lewis
structures
for
the
following
molecules
 –  H2O
 –  NH3
 –  NH4+
 –  CH4
 –  CO32‐
 –  F2
 –  O2
 –  BH3
 –  PCl5
 hypervalent
 –  H2SO4
 –  XeO2
 –  XeF2
 –  XeF4
 –  SF6
 hypervalent
 Lewis
structures
 Resonance
Structures
 ...
View Full Document

This note was uploaded on 10/03/2011 for the course CHEM 113A taught by Professor Professornotknown during the Spring '09 term at San Jose State University .

Ask a homework question - tutors are online